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K. 0. HARTMAN AND I. C. HISATSUNE
dicates a Q value of approximately 1 for ozone formation based on this interpretation of the absorption band. Our current investigations of the spectra and kinetics in the far-ultraviolet have shown at least one other band with a maximum at about 2250 A. Preliminary observations indicate different lifetimes for these bands which must therefore be attributed to different species. The investigations in this spectral region are being continued. Acknowledgment. We are grateful to Mr. E. G. Wendell whose operation and maintenance of the ac-
celerator and assistance with the electronic equipment is vital to these investigations. Miss R. Casey’s contribution to the investigation was greatly appreciated. We are indebted to Professor R. F. Firestone for the use of the analog computer and to Mr. W. Bishop for his helpful advice. Mr. R. Sadler of Argonne National Laboratory kindly provided us with the partially reflecting mirrors. The assistance of Battelle Memorial Institute, Pacific Northwest Laboratories, in supporting W. D. F. as a visiting scientist from Battelle to Ohio State University is gratefully acknowledged.
The Kinetics of Oxalate Ion Pyrolysis in a Potassium Bromide Matrix192a
by K. 0. Hartman2band I. C. Hisatsune Department of Chemistry, Whitmore Laboratory, The Pennsylvania State University, University Park, Pennsylvania 16809 (Received August 16, 1966)
The thermal decomposition of the oxalate ion dispersed in a KBr matrix was observed to follow first-order kinetics with a rate constant of 0.7 X l O I 3 exp[(-60,000 6000)/RT] sec-1. The principal reaction product was the carbonate ion, but traces of formate and monomeric bicarbonate ions were detected as by-products. The yields of the by-products were significant when the initial solute concentrations were low, but at higher concentrations the carbonate ion yield was 90 10%. A reaction mechanism in which the oxalate dissociates into two COz- radicals is proposed.
*
*
Introduction The thermal decompositions of relatively complex oxalate salts have been studied widely, but those of sodium and potassium salts have received much less attention. For example, Glasner and Steinberg3 have carried out extensive studies of the decompositions of rare earth oxalates, and Yankwich and Zavitsanos4 have reported similar studies on magnesium, manganese, and zinc oxalates. Apparently, the only investigation of the kinetics of the pyrolysis of alkali metal oxalates is that reported by Akalan.5 In the present study the infrared disk technique6 was used to follow the decomposition kinetics of the oxalate ion. Our results show that the pyrolysis kinetics in the KBr matrix are different from those in the alkali earth The Journal o j PhysicaE Chemistry
or rare earth oxalate environment reported by earlier investigators. On the basis of kinetic parameters determined in this study, a mechanism for the oxalate ion decomposition is proposed. (1) This work was supported by the National Science Foundation, Grant NSF-G17346, and by the Directorate of Chemical Sciences, Air Force Office of Scientific Research, Grant AF-AFOSR-907-65. (2) (a) Abstracted in part from the Ph.D. Thesis of K. 0. Hartman; (b) Mellon Institute, 4400 Fifth Avenue, Pittsburgh, Pa. 15123. (3) A. Glasner and hl. Steinberg, J . Inorg. AVucl. Chem., 22, 39 (1961); 26, 525 (1964). (4) P. E. Yankwich and P. D. Zavitsanos, J . Phys. Chem., 69, 442 (1965). (5) S. Akalan, Reu. Fac. Sci. Univ. Istanbul, 21, 184 (1956); Chem. Abstr., 51, 9273 (1957). (6) K. 0. Hartman and I. C. Hisatsune, J . Phys. Chem., 69, 583 (1965); 70, 1281 (1966).
KINETICSOF OXALATE ION PYROLYSIS
Experimental Section Kinds and sources of matrix salts used in the present study have been described before.6 Sodium oxalate, a Baker and Adamson product, was recrystallized from water, but potassium oxalate monohydrate from Matheson Coleman and Bell was used directly. The procedures for preparing the pressed disks, for obtaining the kinetic data, and for determining the reaction stoichiometry also have been described before.6 The apparatus for disk fabrication, heating ovens, and infrared instruments were the same as those used in earlier studies. Most of the spectra were obtained on the Perkin-Elmer Model 21 spectrometer.
Results Disks were prepared by both the freeze-dry and grinding methods. I n the freeze-dried samples, the oxalate infrared bands were much sharper than those in the ground samples. As shown in Table I, these two preparative methods resulted in slight frequency variations. The initial spectra of sodium oxalate dispersed in different matrix salts by grinding were essentially the same. Some frequencies varied by a few wavenumbers, but this deviation was within experimental error. When the disks were heated a short time at 500" frequency shifts were observed in KBr, KC1, and KI, but not in NaBr, CsBr, and CsI (Table 11). In those matrices where the shift was observed both sodium and pot:tssium oxalates gave the same set of frequencies after heating. Also, the doublets observed at approximately 1330 and 775 cm-' in the ground samples coalesced into singlets upon heating. Further heating led to the decomposition of the oxalate ion. During the decomposition, a number of bands in addition to those of the major product, carbonate, were generated. From the correlations of the optical densities of these additional bands and from their behavior when the disks were ground or freezedried, the extra bands were found to fall into two sets as shown in Table 111. Since these frequencies agreed well with those of the monomeric bicarbonate ion7 and distorted formate ion,6 the identification of the two byproducts of the reaction was unequivocal. These byproducts were demonstrated, in the following way, to originate from the reaction of the oxalate ion with traces of water present in the disk. A potassium oxalate sample was freeze-dried from a DzO solution and heated at 510' for 1 min. The by-products in this case were identified as bicarbonate4 and formate-d from their characteristic infrared spectra. The quantity of formate and bicarbonate ions produced in different disks was approximately constant and appeared to be independent of the initial oxalate con-
393
Table I : Sodium and Potassium Oxalate Frequencies (cm-1) in Unheated KBr Matrix" --NazCzOr--
Modeb
Freezedried
Ground
KZCZO4' H20
3410 m 2900 w 1332 w
3420 m 2935 w
3400 s 2900 w
WLZ(CO2)
1639 vs
1647 s
1645 s
1421 w
1410 w
WdCOZ)
1416 m 1400 sh 1335 sh 1329 s
1339 sh 1321 s
1325 sh 1310 s
731 s 774 s (515 w)'
780 s
HzO
W,(COZ)
V,(COZ)
779 s
'
Spectra recorded a t 25", sh = shoulder. a = antisymmetric stretch, @ = bending, u = symmetric stretch, w = out-of-plane wag; assignment from G. M. Begun and W. H. Fletcher, Spectrochim. Acta, 19, 1343 (1963). Spectrum on Perkin-Elmer Model 521.
centration. This fact suggested that the secondary reaction was limited by the amount of water present in the disk. This water was very difficult to eliminate. Even when a freeze-dried powder was heated at 100" under vacuum for 24 hr, cooled to room temperature under a dry nitrogen atomosphere, and transferred to the disk die with a maximum of 30-sec exposure to the laboratory air, this sample still produced a detectable quantity of these by-products. A sequence of spectra obtained from a typical kinetic run is illustrated in Figure 1. The initial spectrum is the upper curve, and the second spectrum was obtained after a short heating. Further heating gave the lower spectrum in which the absorption bands of the primary product, carbonate, and the more prominent bands of the minor products, formate and bicarbonate, are present. When the disks were heated for longer periods of time, the oxalate bands disappeared completely followed by the decay of bicarbonate and formate bands. At the commonly employed reaction temperatures (Table V), it was surprising to observe that the carbonate ion absorption bands also decayed slowly. The carbonate ion half-life was about 700 hr at 508". I n addition to the absorption bands of the principal and minor reaction products, we also observed weak (7) D. L. Bernitt, K. 0. Hartman, and I. C. Hisatsune, J. C h m . Phys., 42, 3553 (1965).
Volume 71, Number 2 January 1967
394
K. 0. HARTMAN AND I. C.HISATSUNE
Table 11: Frequencies of Heated Sodium Oxalate in Alkali Halide Matrices (crn-ly Mode" w,(
co
1)
w,(co:!)
KCl
KBr
KI
NsBr
CsBr
CSI
1580 8
1600 s
1595 s
1640 s 1418 w
1640 s 1420 w
1645 s 1420 w
1308
1316 s
1315 s
1336 s 1320 s
1336 s 1320 s
1340 s 1320 s
758 s
760 s
780 s 773 s
780 s 772 s
780 s 773 s
755 s a
Spectra recorded at 25'.
'See footnote b of Table I.
1
Table I11 : Frequencies (cm-l) of By-products Produced in Oxalate Decomposition in KBr Matrix" o
Species 1
Monomeric bicarbonateb
3380 1700 1339 1210 960 835 712
3390 1697 1338 1211 960 835 712
Species 2
Distorted formate'
2660 1630 1445 1350 753
2666 1633 1445 1352 752
11111I
I
I
I
I
I
I
I
I-
1
c
e
x-
n
11
L 3000
Frequency (cm-f)
' Reference 7.
' Spectra recorded a t -190'.
Reference 6.
bands at 2170 and 2331 cm-I which are due, respectively, to the cyanate ion and to carbon dioxide. Trace amounts of cyanate ion were observed in the pyrolyses of nearly all common solid reagents as we reported earlier.' The origin of the carbon dioxide in the matrix is the disproportionation reaction of the carbon monoxide evolved by the oxalate decomposition.6 Our disks, after completion of the reaction, were generally dark gray in color. The stoichiometry of the reaction was determined by comparing the final absorbance of the carbonate bands with calibration curves constructed from disks containing known amounts of carbonate. From more than 20 runs in which the initial oxalate concentration was varied from 0.4 to 2.1 mg/g (milligrams of oxalate per gram of matrix salt), the carbonate yield was 90% f10. Hence, the primary reaction is czo42- =
1
c o p + co
The less than 100% yield of carbonate may in part be accounted for by the slow decomposition of the potassium carbonate as described earlier. The major cause, however, was the formation of the by-products, formate and bicarbonate. It was observed that the yield of carbonate markedly declined when the initial oxalate concentration was decreased below about 0.8 mg/g. The Journal of Physical Chemistry
Figure 1. Infrared spectrum of sodium oxalate in KBr matrix: concentration, 1.5 mg/g; upper spectrum, freeze-dry sample before heating; middle spectrum, after Bmin heating a t 508'; lower spectrum, after 100-min heating a t 508'. All spectra recorded a t 25".
The yield was only 50% when the initial oxalate concentration was 0.403 mg/g and 25% when it was 0.313 mg/g. This reduction in yield was due to the consumption of a higher percentage of the oxalate in the reaction with water to produce the by-products. The reaction order was determined graphically from plots of log rate us. log concentration for the reactant and product absorption bands. The order was found to be 1.0 h 0.2 in oxalate ion. The occurrence of the side reaction with water caused some difficulty in the determination of consistent rate constants for the oxalate decomposition. I n Table IV the variation of the first-order rate constants with initial oxalate concentrations is shown. At lower concentrations where most of the oxalate was consumed by the reaction with water, the rate constant for the loss of oxalate is larger. At such concentrations, the rate of formate and bicarbonate formation was about equal to the rate of oxalate decomposition. Since the amounts of formate and bicarbonate are relatively smaller in higher concentration samples, the rate const,ants from such disks should represent the correct constants for the primary reaction. This expectation is
KINETICS OF
395
OXaLATE ION PYROLYSIS
in fact borne out by the values shown in Table IV. The rate constants are more nearly invariant a t higher concentrations and are equal to the constants determined from the carbonate bands.
Table V: Average Rate Constants for Oxalate Ion Decomposition in KBr Matrix
490 It 5 508 f 8 518 f 8 566 i 8 582 f 7
6.8 18.5 29.0 220.0 405.0
5,7 13.0 13.5 150.0 350.0
4.5 16.0 19.0 175,O 290.0
-Rate 1600 om-1
1.50 1.50 1.51 1.05 0.866 0.866 0.403 0.403 0.313 0.313
1.8 2.0 1.2 1.8 2.2 3.0 3.2 3.0 6.3 6.2
constants, l o 4 k, 8 e c - h 1316 cm-1 758 om-'
1.7 1.8 1.3 2.0 2.3 2.5 3.8 3.0 6.3 6.2
1.5 1.6 1.6 1.8 1.7 2.5 2.3 2.5 6.3 6.2
Activation energy, kcal/mole Frequency factor, sec-l
60 f 6
63 =k 6
60 =k 6
0.91 X 1013
4 . 3 X 1013
0.68 x 1013
COa2-
758 om-' C20r2-
-
QC
Table IV: Effect of Oxalate Concentration on Rate Constants (KBr Matrix)" Initial oxalate concn, mdg
lO'k, sec-1 880 cm-1
1430 cm -1 Coal-
Temp,
To maintain a check on the internal consistency of the results, rare constants were determined throughout the work from both the oxalate 758-cm-' band and the carbonate bands at 1430 and 880 cm-l. The rate constants for the 1430-cm-' carbonate band were higher than those for the other bands, but the discrepancy was within experimental error. Average rate constants determined from these bands are summarized in Table V. The data obtained in the present study are less precise than those obtained in earlier investigationsB because the reaction temperature was not closely controlled. Also, in some runs the 880-cm-' carbonate band split into il doublet with peaks at 883 and 878 cm-l, and this introduced some error in the rate constants. The Arrhenius activation energies and frequency factors determined from different bands are also listed in Table V. For matrices other than KBr, rate constants were determined from single runs at 508". The rate constant obtained using the CsI matrix was the same as that from KBr, while the constants from CsBr and NaBr were about 40% less than this value. Since these runs extended only to 60% reaction this deviation may well be within experimental error.
aL8 The shifts in these frequencies produced by short heating of the disk indicate that the oxalate ion has diffused into the matrix salt and that its environment is now the KBr lattice. Additional evidences which support this interpretation are: the same set of final frequencies was obtained by starting with potassium or sodium oxalate, the frequency shifts were dependent on the matrix salt, and no frequency shifts were observed when undiluted oxalate salts were heated. We may also conclude from the fact that shifts were not observed in NaBr, CsBr, and CsI matrices that the extent of diffusion of the solute ion into these matrices is negligible. The kinetic data obtained in this study thus can be associated with the decomposition of oxalate ions isolated in the KBr matrix. The observed first-order rate law, the experimental activation energy, and the calculated frequency factor all suggest that this decomposition can be regarded as a unimolecular reaction. A simple mechanism which accounts for the experimental data is one in which the rate-determining step is the unimolecular dissociation of the oxalate ion into two carbon dioxide anion radicals. This step is followed by a rapid bimolecular reaction of the two free radicals to form the products carbonate ion and carbon monoxide. A possible transition complex for the second step may be the carbonyl carbonate ion, OCOCOZ2- which is isoelectronic with a known isomer of Kz04.9 The formation of bicarbonate and formate as by-products of the decomposition reaction is not inconsist.ent with this free-radical mechanism since the COz- radical has been observed to react readily with water to form just these
Discussion The frequencies of the sodium oxalate absorption bands observed in the KBr matrix before heating are in reasonable agreement with those reported by Schmelz,
(8) M. J. Schmelz, T. Miyazawit, 5. hliaushima, T. J. Lane, and J. V. Quagliano, Spectrochim. Acta, 9,51 (1957). (9) I. C. Hisatsune, J. P. Devlin, and Y . Wada, J . Chem. Phys., 3 3 , 714 (1960).
All samples were prepared by the freeze-dry method. action temperature = 508'.
et
Re-
Volume 71, Number 2 January 1967
396
products.1° However, our experimental data do not allow us to determine whether these by-products were formed by a concerted attack on the oxalate ion by water or through the direct reaction of the Con- radicals with water. We also examined several disks for the presence of the COz- radical but found no esr signal that can be attributed to this species. This negative result, however, does not contradict our mechanism because previously we found the COZ- radical to be unstable a t t,emperatures above about 100" in alkali halide matrices.'O The only esr signal observed in our pressed disks was a sharp line with g = 2.0028. The same signal was observed previously in our formate kinetic studies.6 This free radical appears to be an ion containing only oxygen atom or atoms since no hyperfine structures were observed with H, D, and 13C isotopic samples. The observed Arrhenius activation energy of 60 f 6 kcal/mole is considerably higher than 32 and 43 kcal/ mole reported by Akalan; respectively, for the sodium
The Journd of Physical Chemktry
K. 0. HARTMAN AND I. C. HISATSUNE
and potassium salts. However, our present value is in agreement within experimental error with the 68 f 5 kcal/mole determined for calcium oxalate pyrolysis by both the infrared and thermogravimetric methods. l1 Also, the calculated rate constant of 3.2 X min-' at 450" is in reasonable agreement with 1.7 X min-l found recently by Wing and Harris12 for the decomposition of undiluted potassium oxalate. In addition, the reaction mechanism proposed by the latter authors involved a step similar to our rate-determining dissociation of an oxalate ion into two COZfree radicals.
Acknowledgments. We are pleased to acknowledge the financial support from the National Science Foundation and the Air Force Office of Scientific Research. (10) IC. 0.Hartman and I. C. Hitsatsune, J. Chem.Phys., 44, 1913 (1966). (11) F. E. Freeberg, K. 0. Hartman, I. C. Hisatsune, and J. W. Schempf, J . Phys. C h a . , 71, 397 (1967). (12) R. M.Wing and G. M. Harris, ibid., 69,4328 (1965).
